Download 3.1.2 Group 2

3.1.2 Group 2
Atomic radius
Atomic radius increases down the Group.
As one goes down the group the atoms have more
shells of electrons making the atom bigger
Electronic Structure
Group 2 metals all have the outer shell s2 electron
configuration.
Melting points
Down the group the melting points decrease. The
metallic bonding weakens as the atomic size
increases. The distance between the positive ions and
delocalized electrons increases. Therefore the
attractive forces between the positive ions and the
delocalized electrons weaken.
Ionisation energy
When the group 2 metals react they lose the outer shell s2 electrons in redox reactions to form 2+ ions. The
energy to remove these electrons are the first and second ionisation energy
The first ionisation energy is Energy needed to remove an
electron from each atom in one mole of gaseous atoms
This is represented by the equation:
H(g)
H+(g) + e-
Always gaseous
Second ionisation energy
The second ionisation energy is the enthalpy change when one mole of
gaseous ions with a single positive charge forms one mole of gaseous
ions with a double positive charge
This is represented by the equation:
Group 2 reactions
Ti+ (g)
The first and second ionisation
energies decrease down the group.
The outermost electrons are held more
weakly because they are successively
further from the nucleus in additional
shells
In addition, the outer shell electrons
become more shielded from the
attraction of the nucleus by the
repulsive force of inner shell electrons
Ti2+(g) + e-
Reactivity of group 2 metals increases down the group
The reactivity increases down the group as the atomic radii increase there is more shielding.
The nuclear attraction decreases and it is easier to remove (outer) electrons and so cations
form more easily
Reactions with oxygen.
The group 2 metals will burn in oxygen.
Mg burns with a bright white flame
2MgO
2Mg + O2
MgO is a white solid with a high melting
point due to its ionic bonding
Mg will also react slowly with oxygen without a flame.
Mg ribbon will often have a thin layer of magnesium oxide on it
formed by reaction with oxygen
2Mg + O2
2MgO
This needs to be cleaned off by emery paper before doing
reactions with Mg ribbon
If testing for reaction rates with Mg and acid, an un-cleaned Mg
ribbon would give a false result because both the Mg and MgO
would react but at different rates.
Mg + 2HCl
MgCl2 + H2
MgO + 2HCl
MgCl2 + H2O
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Reactions with water.
Magnesium burns in steam to produce
magnesium oxide and hydrogen. The Mg
would burn with a bright white flame
Mg (s) + H2O (g)
MgO (s) + H2 (g)
Mg will also react with warm water, giving a different
magnesium hydroxide product
Mg + 2 H2O
Mg(OH)2 + H2
This is a much slower reaction than the reaction with
steam and there is no flame
The other group 2 metals will react with cold water with
increasing vigour down the group to form hydroxides
The hydroxides produced make the water alkaline
One would observe:
Ca + 2 H2O (l)
Ca(OH)2 (aq) + H2 (g)
•fizzing, (more vigorous down group)
Sr + 2 H2O (l)
Sr(OH)2 (aq) + H2 (g)
•the metal dissolving, (faster down group)
Ba + 2 H2O (l)
Ba(OH)2 (aq) + H2 (g)
•the solution heating up (more down group)
•and with calcium a white precipitate
appearing (less precipitate forms down
group)
Reactions with acid
The group 2 metals will react with acids with increasing
vigour down the group to form a salt and hydrogen
Ca + 2HCl (aq)
CaCl2 (aq) + H2 (g)
Sr + 2 HNO3 (aq)
Sr(NO3)2 (aq) + H2 (g)
Mg + H2SO4 (aq)
MgSO4 (aq) + H2 (g)
If Barium metal is reacted with sulphuric acid it
will only react slowly as the insoluble Barium
sulphate produced will cover the surface of the
metal and act as a barrier to further attack.
Ba + H2SO4
BaSO4 + H2
The same effect will happen to a lesser extent
with metals going up the group as the solubility
increases.
The same effect does not happen with other
acids like hydrochloric or nitric as they form
soluble group 2 salts.
Action of water on oxides of elements in Group 2
The group 2 oxides react with water to form hydroxides of
varying solubility
CaO (s) + H2O (l)
Ca(OH)2 (aq) pH 12
MgO (s) + H2O (l)
Mg(OH)2 (s) pH 9
Mg(OH)2 is only slightly soluble in water so fewer free
OH- ions are produced and so lower pH
Magnesium hydroxide is classed as partially soluble in
water.
A suspension of magnesium hydroxide in water will appear
slightly alkaline (pH 9) so some hydroxide ions must
therefore have been produced by a very slight dissolving.
Magnesium hydroxide is used in medicine (in suspension
as milk of magnesia) to neutralise excess acid in the
stomach and to treat constipation.
Mg(OH)2 + 2HCl
Group 2 oxides are basic as the oxide ions
accept H+ ions to become hydroxide ions in
these reactions
Calcium hydroxide is reasonably soluble in
water. It is used in agriculture to neutralise
acidic soils.
If too much calcium hydroxide is added to the
soil, excess will result in soils becoming too
alkaline to sustain crop growth
An aqueous solution of calcium hydroxide is
called lime water and can be used a test for
carbon dioxide. The limewater turns cloudy as
white calcium carbonate is produced.
Ca(OH)2 (aq) + CO2 (g)
MgCl2 + 2H2O
It is safe to use as it so weakly alkaline.
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CaCO3 (s) + H2O(l)